In previous articles, we have seen formation of ionic bonds. In this article, we shall study the concept of the formation of a covalent bond.

Covalent Bonds or Electron Pair Bonds:

A bond established between two identical or different atoms by sharing one or more pairs of electrons is known as a covalent bond. The attractive force which comes into existence due to the mutual sharing of electrons between two atoms having similar electronegativity or having a small difference in electronegativities is called a covalent bond.

No ionic bond is possible between two atoms having similar electronegativities. Hence to explain the bonding between such atoms Lewis introduced the concept of covalent bonds.

Each atom contributes one electron to form a common pair i.e. equal contribution of electrons followed by equal sharing. If one electron pair is shared, it is known as a single covalent bond.  If two electron pairs are shared, it is a double covalent bond and so on. In CI2, O2, CH4, H2O, etc. there are covalent bonds.

Examples of Covalent Bonds:

Formation of Hydrogen Molecule:

Electronic configuration of hydrogen H (Z = 1) is (1) Hydrogen has one electron in its valence shell. It tries to acquire a stable configuration (2) by sharing its one electron with another hydrogen atom.

When two hydrogen atoms approach each other, at a certain distance between the nuclei, they share their valence electrons and form a shared pair of electrons. The shared pair belongs to each atom equally. As only one pair is shared between the two hydrogen atoms, they are joined to each other by a single covalent bond.

Formation of Chlorine Molecule:

Electronic configuration of chlorine Cl (Z = 17) is (2, 8, 7). Chlorine has 7 electrons in its valence shell. It can acquire a stable configuration (2, 8, 8) of the nearest inert gas (Ar) by sharing one electron with another chlorine atom.

When two chlorine atoms approach each other, at a certain distance between the nuclei, they share their valence electrons and form a shared pair of electrons. The shared pair belongs to each atom equally. As only one pair is shared between the two chlorine atoms, they are joined to each other by a single covalent bond.

Formation of Hydrogen Chloride Molecule:

Electronic configuration of hydrogen H (Z = 1) is  (1)  Hydrogen has one electron in its valence shell. It tries to acquire a stable configuration (2) by sharing its one electron with the chlorine atom. Electronic configuration of chlorine Cl (Z = 17) is (2, 8, 7). Chlorine has 7 electrons in its valence shell. It can acquire a stable configuration (2, 8, 8) of nearest inert gas (Ar) by sharing one electron with the hydrogen atom.

When hydrogen and chlorine atoms approach each other, at a certain distance between the nuclei, they share their valence electrons and form a shared pair of electrons. The shared pair belongs to each atom equally. As only one pair is shared between the hydrogen and chlorine atoms, they are joined to each other by a single covalent bond.

Formation of Ammonia Molecule:

Electronic configuration of nitrogen N (Z = 7) is (2, 5). Nitrogen has 5 electrons in its valence shell. It can acquire stable configuration (2, 8) of nearest inert gas (Ne) by sharing one electron each with three hydrogen atom.

When three hydrogen atoms approach nitrogen atom, at a certain distance between the nuclei, it shares one valence electron each with the nitrogen atom and forms three shared pairs of electrons with three hydrogen atoms. The shared pair belongs to each atom equally. As only one pair is shared between the atoms, they are joined to each other by a single covalent bond.

Formation of Water Molecule:

Electronic configuration of oxygen O (Z = 8) is (2, 6). Oxygen has 6 electrons in its valence shell. It can acquire stable configuration (2, 8) of nearest inert gas (Ne) by sharing one electron each with two hydrogen atom.

When two hydrogen atoms approach oxygen atom, at a certain distance between the nuclei, it shares one valence electron each with two hydrogen atoms and forms two shared pairs of electrons with two hydrogen atoms. The shared pair belongs to each atom equally. As only one pair is shared between the atoms, they are joined to each other by a single covalent bond.

Formation of Phosphorous Trichloride Molecule:

Electronic configuration of phosphorous P (Z = 15) is (2, 8, 5). Phosphorous has 5 electrons in its valence shell. It can acquire stable configuration (2, 8, 8) of nearest inert gas (Ar) by sharing one electron each with three chlorine atoms. Electronic configuration of chlorine Cl (Z = 17) is (2, 8, 7). Chlorine has 7 electrons in its valence shell. It can acquire a stable configuration (2, 8, 8) of the nearest inert gas (Ar) by sharing one electron with another atom.

When three chlorine atoms approach phosphorous atom, at a certain distance between the nuclei, it shares one valence electron each with the three chlorine atoms and forms three shared pairs of electrons with three chlorine atoms. The shared pair belongs to each atom equally. As only one pair is shared between the atoms, they are joined to each other by a single covalent bond.

Formation of Methane Molecule:

Electronic configuration of carbon C (Z = 6) is (2, 4). Carbon has 4 electrons in its valence shell. It can acquire stable configuration (2, 8) of nearest inert gas (Ne) by sharing one electron each with four hydrogen atoms.

When four hydrogen atoms approach phosphorous atom, at a certain distance between the nuclei, it shares one valence electron each with the four hydrogen atoms and forms four shared pairs of electrons with four hydrogen atoms. The shared pair belongs to each atom equally. As only one pair is shared between the atoms, they are joined to each other by a single covalent bond.

Formation of Ethane Molecule:

Formation of Oxygen Molecule:

Consider the formation of the oxygen molecule. Electronic configuration of oxygen O (Z = 8) is (2, 6). Oxygen has 6 electrons in its valence shell. It can acquire a stable configuration (2, 8) of nearest inert gas (Ne) by sharing two electrons with another oxygen atom. The two atoms share two pairs of electrons between them completing octet of each. two shared pairs constitute a double bond.

O        +      O        →   O=O

(2,6)      (2,6)       (2,8) (2,8)

Formation of Nitrogen Molecule:

Consider the formation of the nitrogen molecule. Electronic configuration of nitrogen (Z = 7) is (2, 5). Nitrogen has 5 electrons in its valence shell. It can acquire a stable configuration (2, 8) of nearest inert gas (Ne) by sharing three electrons with another nitrogen atom. The two atoms share three pairs of electrons between them completing octet of each. Three shared pairs constitute a triple bond.

 N        +      N        →   N≡N

(2,5)      (2,5)       (2,8) (2,8)

Formation of Ethene (Ethylene) Molecule:

Formation of Ethyne (Acetylene) Molecule:

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > Formation of Covalent Bonds

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In this article, we shall study the properties of ionic compounds.

Physical State:

Due to strong electrostatic force present between the oppositely charged ions, they are held closer and fixed at specified positions in the crystal lattice. Hence ionic compounds usually exist in the form of crystalline solids at room temperature. Actually, these compounds do not possess a molecule but it is a well-defined geometric arrangement of positive and negative ions in the crystal lattice.

In sodium chloride (NaCl), each Na⁺ ion is surrounded by 6 Cl^– ions and each Cl^– ion is surrounded by 6 Na⁺ ions. The arrangement is as shown below.

High Melting and Boiling Points of Ionic Compounds:

Due to strong electrostatic force present between the oppositely charged ions, they are held closer and fixed at specified positions in the crystal lattice. Hence very high temperature is required to separate from each other and to move freely as in the case of liquid. Hence the boiling and melting points of ionic compounds are generally very high.

Hard and Brittle Nature of Ionic Compounds:

Due to strong electrostatic force present between the oppositely charged ions, they are held closer and fixed at specified positions in the crystal lattice. There are layers of positive ions and negative ions. When force is applied to them the layer may slide and the ions having similar charges may come near each other. In such a case, the two layers repel each other and the crystal gets cleaved. Hence ionic solids are brittle.

Electrical Conductivity:

Due to strong electrostatic force present between the oppositely charged ions, they are held closer and fixed at specified positions in the crystal lattice. Hence in solid-state ionic compounds do not conduct electricity. In the Fused state or the dissolved state, the ions of the ionic compound can move from one position to another freely. Hence in fused state, ionic compounds conduct electricity.

Lattice Energy (Δ_latticeH):

The enthalpy change involved in the formation of one mole of an ionic crystal from its constituent gaseous positive and negative ion is called lattice energy. The formation of an ionic compound is an exothermic process. Hence the enthalpy change involved in the process of formation of an ionic compound is negative. The lattice energy decreases with the increase in the ionic size.

Hydration Energy (Δ_hydrationH|):

The enthalpy change involved in the hydration of one mole of gaseous ions of each type of an ionic compound (solid) is called hydration energy. Hydration of ionic compounds is exothermic. Hence the enthalpy change involved in the process of hydration of an ionic compound is negative. The hydration energy decreases with the increase in the ionic size.

Solubility in water:

Water has a high dielectric constant hence it weakens the electrostatic force between the ions of an ionic compound. Due to which the ions get separated and get surrounded by the water molecule. This dissolves the ionic compound in water. Hence ionic compounds are soluble in water.

Organic solvents like benzene, ether, hexane have low dielectric constant and are unable to weaken the electrostatic force between the ions of the ionic compound. Thus they are unable to separate the ions. Hence ionic compounds are not soluble in organic solvents.

The enthalpy of solution of an ionic solid is numerically equal to the difference in its hydration energy and lattice energy

Mathematically,   Δ_solutionH  = Δ_hydrationH – Δ_latticeH

The Gibb’s energy (free energy) change is given by

ΔG  = ΔH – T ΔS    i.e.

Δ_solutionG  = Δ_solutionH – T Δ_solutionS

Hence for a compound to get dissolved in water, Gibb’s energy (free energy) change involved in the solution formation process is negative. Thus the quantity (Δ_solutionH – T Δ_solutionS) is negative. The quantity Δ_solutionS is positive i.e. the quantity T Δ_solutionS is positive. It means the quantity Δ_solutionH should be smaller than T Δ_solutionS. Hence a more negative value of Δ_solutionH will help the greater dissolution of an ionic compound in water.

Hence we can conclude that

• If Δ_hydrationH > Δ_latticeH, the salt would dissolve in water.
• If Δ_hydrationH < Δ_latticeH, ordinarily the salt would dissolve in water. It will start dissolving when Δ_solutionH < T Δ_solutionS

Stereoisomerism:

Isomers having the same molecular formula, same structural formula but different configurations are called stereo Isomers and the phenomenon is known as stereoisomerism. The ionic compounds are neither rigid or directional hence they do not show stereoisomerism.

Ionic Reactions:

The ionic reaction involves the mutual interaction of ions. Due to the presence of ions,  the reactions of ionic compounds are very fast

Born-Haber Cycle

The energy change involved in the formation of ionic bonds from the constituent elements can be represented by a cycle called Born-Haber Cycle. It uses Hess’s law of thermodynamics to calculate the change in enthalpy during ionic bond formation.

Let us consider the formation of 1 mole of sodium chloride from sodium and chlorine.

The Born-Haber Cycle can be simply stated as Heat of formation of ionic crystal = Heat of atomization + Dissociation energy+ (sum of Ionization energies)+ (sum of Electron affinities) +  Lattice energy.

If energy is released (exothermic reaction), put a negative sign in front of the value of the energy; if energy is absorbed (endothermic reaction), the value of energy should be positive.

Note:

In this general equation, the electron affinity is added. However, when plugging in a value, determine whether energy is released (exothermic reaction) or absorbed (endothermic reaction) for each electron affinity.

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > Properties of Ionic Compounds

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In this article, we shall study the factors affecting the formation of the ionic bond and the concept of variable electrovalency.

Ionization energy of Electropositive Atom:

It is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom of an element. The ionization energy of an atom can be considered as a measure to lose an electron and form a cation. The lesser the ionization energy, the greater is the ease of the formation of a cation and thus it can form ionic bond easily.

Alkali metals and alkaline earth metals have low ionization energy hence they have a tendency to form ionic compounds.

Electron Affinity or Electron Gain Enthalpy of Electronegative Atom:

It is defined as the amount of energy released when an electron is added to an isolated gaseous atom of an element. The electron gain enthalpy of an atom can be considered as a measure to gain an electron and form an anion. The higher the energy released during this process, the easier will be the formation of an anion.

Elements of groups 16 and 17 have more negative values of electron gain enthalpies hence they have a tendency to form ionic compounds. 

Thus, the low ionization energy of a metal atom and high electron affinity of a non-metal atom facilitate the formation of an ionic bond between them.

Lattice energy or Lattice Enthalpy:

The enthalpy change involved in the formation of one mole of an ionic crystal from its constituent gaseous positive and negative ion is called lattice energy. The formation of an ionic compound is an exothermic process. Hence the enthalpy change involved in the process of formation of an ionic compound is negative.

Higher the lattice energy, the greater is the tendency of the formation of an ionic bond. The higher the charges on the ions and the smaller the distance between them, the greater is the force of attraction between them.

The lattice energy depends on the following factors

• Size of the atom: If the ions formed are smaller in size, the inter-nuclear distance is less and hence the inter-ionic attractive force is greater. Hence the lattice energy will be high. Thus smaller the size of ions, greater is the lattice energy.
• Charge on the ions: Higher the charges on the ion, greater the inter-ionic attractive force. Hence lattice energy is high. Thus higher the charge on the ions, greater is the lattice energy.

Difference in Electronegativities:

The tendency of an atom to attract the bonding or a shared pair of electrons towards its own side in a covalent bond is called electronegativity of that atom. Higher the difference in electronegativities of the two atoms, greater will be the ease to form an ionic bond.

A difference of 2 units (pauling) in electronegativities is required to form an ionic bond.

The Concept of Variable Electrovalency:

There are many elements in the periodic table which are capable of forming more than one type of ions having different charges. Thus they possess more than one electrovalency. This phenomenon is known as a variable electrovalency. For example, copper forms cuprous (Cu⁺) and cupric (Cu ²⁺) ions. Iron forms ferrous (Fe²⁺) and ferric (Fe³⁺) ions.

Causes of Variable Electrovalency:

Unstable Nature of Core (Kernel):

When an atom loses one or more electrons, a cation is formed. The remaining part of the atom left is called a core or a kernel of the atom.

During the formation of positive ion-neutral atom loses one or more (definite) number of electrons to form a cation. If core or kernel form is stable then it will exhibit definite electrovalency.

Na               →      Na⁺       +     e^–

1s² 2s² 2p⁶ 3s¹          1s² 2s² 2p⁶

Sodium atom   core or kernel of sodium

If core or kernel of ion formed is not stable then to acquire greater stability it is compelled to lose more electrons. This fact gives rise to variable electrovalency. Let us consider the case of iron Fe (Z = 26)

Fe                   →          Fe ²⁺       +    2 e^–

1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²       1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶

(2, 8, 14, 2)                          (2, 8, 14)  less stable

Fe²⁺               →      Fe ³⁺       +     e^–

1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶        1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵

(2, 8, 14)                          (2, 8, 13)

more stable  due to exactly half filled d orbitals

Inert Electron Pair Effect:

The reluctance of ns² electron pair to get excited and to take part in bond formation is called inert electron pair effect.

Let us consider case of tin Sn (Z = 50). Electronic configuration of tin is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s² 5p². (2, 8, 18, 18, 4). It has 4 electrons and can form Sn⁴⁺ ion by donating 4 electrons and acquires configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ (2, 8, 18, 18).

But due to reluctance of 5s² electron pair to get excited and to take part in bond formation, it loses only two electrons and acquires configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ (2, 8, 18, 2) an forms Sn²⁺ ion.

Thus tin shows two electrovalencies due to the inert pair effect.

Science > Chemistry > Physical Chemistry > Nature of Chemical Bond > Factors Governing Formation of Ionic Bond

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The force of attraction which keeps atoms or ions together in a molecule is called a Chemical bond. The main cause of chemical combination is the tendency to acquire stability i.e. a state with minimum energy and a tendency to acquire noble gas configuration in outermost orbit. Many theories are given to explain the nature and types of chemical bonds. In this article, we shall study the octet theory and different types of bonds based on it.

Types of Chemical Bond:

• Electrovalent or ionic bond
• Covalent bond
• Coordinate or Dative bond
• Metallic bond

Octet theory or Octet Rule:

It was put forward by Kossel and Lewis (1916).  It states that

• Electronic configuration of noble gases is most stable. Except for helium, all other noble gases have eight electrons in the outermost shell.
• Atoms combine with each other by loss, gain or sharing of electrons so as to acquire an octet of electrons (i.e. 8 electrons) in their outermost shell to attain configuration of nearest noble (inert)

Lewis Structure:

Electrons in inner orbit do not take part in a chemical reaction. For forming bonds only electrons in outermost orbit are involved. The electrons present in the outermost shell of an atom are called valence electrons. Lewis introduces a simple notation to represent these valence electrons by dots. Only valence electrons are shown as dots surrounding the symbol of the given element.

Classification of Chemical Bonds on the Basis of Octet Theory:

According to octet theory, chemical bonds are classified into three types as :

• Ionic or Electrovalent bond.
• Covalent bond.
• Coordinate bond.

Ionic or Electrovalent Bond:

The electrostatic force of attraction between the oppositely charged ions is called the ionic bond or electrovalent bond. In this type of bond formation, there is a transfer of electrons from the outermost shell of one atom (generally strongly electropositive) to the outermost shell of the other atom (generally strongly electronegative).

Formation of Sodium Chloride:

In the formation of sodium chloride, an electron is transferred from sodium atom to chlorine atom and a positive ion of sodium and negative ion of chlorine is formed. The resulting Na⁺ and Cl^– ions, possessing configuration of neon (2,8) and argon (2,8,8) respectively combine to form an electrovalent compound as shown below:

(2,8,1)               (2,8,7)               (2,8)               (2,8,8)

Formation of Magnesium Chloride:

In the formation of magnesium chloride, two electrons are transferred from magnesium atom to two chlorine atom (one each) and positive ion of magnesium and two negative ions of chlorine is formed. The resulting Mg⁺⁺ and two Cl^– ions, possessing configuration of neon (2,8) and argon (2,8,8) respectively combine to form an electrovalent compound (MgCl₂) as shown below:

In the above examples, each ion has attained an octet in the outermost shell. When positive and negative ions come closer to each other, they are held by electrostatic forces of attraction.

Covalent Bond:

A bond established between two identical or different atoms by sharing one or more pairs of electrons is known as a covalent bond. Each atom contributes one electron to form a common pair i.e. the equal contribution of electrons followed by equal sharing. If one electron pair is shared, it is known as a single covalent bond.  If two electron pairs are shared, it is a double covalent bond and so on.

It is observed in diatomic molecules like chlorine (CI2), oxygen (O2), hydrogen(H2), nitrogen (N2), fluorine (F2) etc. and other molecules like methane (CH4), water (H2O), carbon dioxide(CO2) etc 

Formation of Chlorine Molecule:

Consider formation of chlorine molecule.  Each atom of chlorine consists of seven valence electrons in outermost shell. The two atoms share one pair of electrons between them completing octet of each. One shared pair constitutes a single bond.

Cl  +  Cl   →   Cl – Cl i.e. Cl₂

(2,8,7) (2,8,7)  (2,8,8) (2,8,8)

           Chlorine atoms        Shared pair

Formation of Oxygen Molecule:

1. Consider formation of oxygen molecule.  Each atom of oxygen consists of six valence electrons in outermost shell. The two atoms share two pairs of electrons between them completing octet of each. two shared pairs constitutes a double bond.

O  +  O   →   O = O i.e. O₂

(2,6) (2,6)  (2,8) (2,8)

           Oxygen atoms        Shared two pairs

Formation of Nitrogen Molecule:

Consider formation of nitrogen molecule.  Each atom of nitrogen consists of five valence electrons in outermost shell. The two atoms share three pairs of electrons between them completing octet of each. three shared pairs constitutes a triple bond.

N  +  N   →   N ≡ N i.e. N₂

(2,5) (2,5)  (2,8) (2,8)

           Nitrogen atoms        Shared three pairs

Coordinate Bond:

Coordinate bond is a special type of covalent bond When one atom has complete octet with the unshared pair of electrons (lone pair) and another atom is in need of a pair of electrons then they Form bond what is known as a coordinate bond. Both the electron required for sharing between two atoms are contributed by one atom only. The atom contributing the lone pair is known at Donor and the atom in need of two electrons is known as Acceptor.

Coordinate bond is represented by → from donor to the acceptor.

In ammonia nitrogen has a complete octet with a lone pair of electrons.  It can donate the lone pair to hydrogen ion forming NH⁴⁺ ion.

Difference Between Covalent Bond and Co-ordinate Bond:

A bond established between two identical or different atoms by sharing of one or more pairs of valence electrons in the outermost shell is known as covalent bond while the coordinate bond is a special type of covalent bond in which only a single atom whose octet is complete contributes both the electrons required for sharing between two atoms.

In the formation of a covalent bond, each atom contributes one electron to form a common pair. i.e. the equal contribution of electrons followed by equal sharing, while in the formation of coordinate bond in which only a single atom whose octet is complete contributes both the electrons required for sharing between two atoms.

Inadequacies or Limitation of Octet Theory: (Need of Valence Bond Theory)

Failure to Explain Incomplete and expanded Octets:

It fails to explain the formation of ions and compounds with incomplete as in BeCI₂ and BCI₃ or expanded octet as in PCI₅ and SF₆   Octet rule in not obeyed in these cases, still, compounds are formed.

In the above structures, neither the octet of Be (4 – electrons around Be) nor of B (6 – electrons around B) is completed.  These atoms have incomplete octets, however, they are stable compounds. The octet theory is unable to explain this phenomenon.

The central atoms phosphorus and sulphur have ten and twelve electrons around them respectively.  Such molecules are said to have expanded octet.  Octet rule in not obeyed in these cases, still compounds are formed.

Failure to explain the Nature of interacting forces:

The octet rule cannot explain the nature of the forces of interaction between two combining atoms.  (There are attractive forces between the nucleus of one atom and electrons of other atom and there are repulsive forces between nucleus-nucleus and electron-electron of  two atoms).

Octet theory fails to explain the nature of interacting forces it can be explained as follows

• the octet theory simply says that a covalent bond will be formed by sharing of electrons but tells nothing about the interacting forces.
• There are attractive forces between the nucleus of one atom and electrons of other atom and there are repulsive forces between nucleus-nucleus and electron-electron of two atoms.
•  Octet theory does not explain what forces make the atoms to combine with each other.

Failure to explain Energy and reactivity:

It does not explain the energy relations of atoms and molecules.  The molecule has lower energy and greater stability than the combining atoms. The molecule has lower energy and greater stability than the combining atoms. Octet theory fails to explain the energy evolved in bond formation and reactivity of molecules formed.

The energy given off or released per mole during the formation of a covalent bond is called bond energy. Greater the bond energy, the more stable will the bond and less reactive will be the molecule. The octet theory does not enable us to calculate the bond energy, in a chemical bond.

Octet theory fails to explain the reactivity of covalent molecules. Once the octet of all involved atoms is completed, the molecule should be as unreactive as an inert gas. Octet theory fails to explain the energy and reactivity of the resulting molecule. It can be explained as follows

Failure to explain the Geometry of the molecules:

It fails to explain the geometry of the molecule.  Molecules have different shapes such as linear, trigonal, tetrahedral etc. Octet theory fails to explain the geometry of the resulting molecule. It can be explained as follows

• Molecules have different shapes such as linear, trigonal, tetrahedral, etc.
• The octet theory provides an idea about the number of covalent bonds formed but not the geometry of the molecules formed.
• Octet theory fails to explain the tetrahedral structure of methane, the distorted tetrahedral structure of ammonia and water, the trigonal planar structure of boron trifluoride and the linear structure of beryllium difluoride.
• BeF2 is a linear molecule while H₂O is a v-shaped molecule. The octet theory fails to explain this difference.

Octet theory is only a guideline for writing “dot and dash” structures

Thus octet theory has a number of limitations and therefore another theory (valence bond theory) is required to explain the nature of the chemical bond.

Dot and Dash Structures Using Octet Theory:

Electrons are generally represented by a dot (.) or cross (x) and a covalent bond is expressed by a dash  (-).  The system of representing the formula of a compound using dot and dashes is called “Electron dot and dash formula”.

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